reaction energy and reaction kinetics general chemistry unit 12
TRANSCRIPT
Reaction Energy and Reaction Kinetics
General ChemistryUnit 12
Driving Forces
• Enthalpy and Entropy• Enthalpy (heat of reaction) is the
amount of energy released or absorbed during a chemical reaction • Symbol is ΔH• Think of it as energy needed
Thermochemical Equations
• A thermochemical equation shows the energy (enthalpy) change in the reaction• Put in as reactant or product
2 H2 + O2 → 2 H2O + 483.6 kJ
• List behind as ΔH2 H2 + O2 → 2 H2O ΔH = -483.6 kJ
• If energy is released (product) the reaction is exothermic and ΔH is negative
• If energy is absorbed (reactant) the reaction is endothermic and ΔH is positive
Entropy
• Entropy is a measure of randomness, tendency toward disorder• Symbol is ΔS• More disorder = more entropy• If reaction leads to more disorder, the
entropy change (ΔS) is positive, if it becomes more ordered, ΔS is negative
• Example: melting ice, condensing water, cleaning your room (+,-,-)
Free Energy (ΔG)• Free energy combines enthalpy and
entropy to measure the spontanaeity of a reaction
• Gibbs Free Energy Equation:ΔG = ΔH - T ΔS (T is in Kelvin: +273 to ºC)
• If ΔG is negative, reaction is spontaneous
• If ΔG is positive, reaction is NOT spontaneous, but would be spontaneous in the reverse direction
ExampleFind ΔG for the reaction:
NH4Cl(s) → NH3(g) + HCl(g)
Using the following data:ΔH = 176 kJ, ΔS = 285 J/K, T = 25ºC
Solution: (Change to kJ and K) ΔG = 176 kJ – (298 K)(.285 kJ/K) ΔG = 176 kJ – 84.9 kJ
= 91 kJ NOT spontaneous
Comparison of Signs
ΔH ΔS ΔGSpontaneous?
- + - ALWAYS spont. + - + NEVER spont.
- - - / + Spont. at low T + + - / + Spont. at high T
Reaction Mechanisms
• Step-by-step sequence that occurs to create the products
• Intermediates may form that do not appear in overall reaction – they are used up in another step
• Homogeneous reaction: all reactants in same phase
• Heterogeneous reaction: reactants in different phases
• Rate-determining step: slowest step of reaction mechanism
Activation Energy
• Minimum energy to make the reaction go (form activated complex which allows reaction to proceed)
• Reaction needs:• Enough energy• Proper orientation of molecules –
must hit each other at correct spot
Energy Diagrams
Exothermic/Endothermic
Energy Example
1. Calculate the ΔH.20 kJ – 40 kJ = -20 kJ
2. Calculate the Ea.
100 kJ – 40 kJ = 60 kJ
3. Calculate the Ea‘.
100 kJ – 20 kJ = 80 kJ
Reaction Rate
• Rate can be defined in terms of molar concentration (M) for the disappearance of a reactant or the appearance of a product
• Concentration shown as:[HCl] = 0.1 means the molar concentration of HCl is 0.1 M
Factors Affecting Reaction Rate
• Nature of reactants• Concentration• Temperature• Catalysts
Nature of Reactants
• Ionic – almost instantaneous• Molecular – slower (bonds must
break and reform)• Surface area – rate increases with
greater surface area
Concentration
• Measured in molarity [A]• Increasing the concentration of
reactants increases the rate• Rate law:
Rate = k[A]m[B]n
The exponents m and n must be determined experimentally
Temperature
• Increasing the temperature gives more collisions between molecules
• This leads to the formation of more activated complexes and this causes the rate to increase
↑ T → ↑ collisions → ↑ complexes → ↑ rate
Catalysts• Catalyst – increase reaction rate without
being used up• Lower the activation energy Animation
• Heterogeneous – not in same phase as reactants, provides surface to give more effective collisions Catalytic Converter
• Homogeneous – In same phase as reactants, makes different activated complex, returns to original form at end of reaction
• Demo: Catalysts